Download Unit 13: Electrochemistry (Link to Prentice Hall Text: Chapters 22

Unit 13: Electrochemistry
(Link to Prentice Hall Text: Chapters 22 & 23)
Name:____________________________________________________________________________________
Date Due
Assignments
Page Number: Problem Numbers
Assignment 1: Balancing and Identifying REDOX Reactions
673: 30, 32, 34
Assignment 2: Electrochemical Cells
701: 22, 34, 35, 38, 42
A. Atoms Compete for Electrons (Electrochemistry)
Defining Oxidation and Reduction
When chemical bonds form, electrons are either lost, gained or shared. In REDOX reactions, electrons are lost or gained.
Oxidation Is Loss, Reduction Is Gain, “OIL RIG”
Oxidation:
-
Loss of electrons.
-
Metals are more easily oxidized.
-
They are termed “reducing agents.”
-
If a substance is oxidized, its oxidation number increases.
Reduction:
-
Gain of electrons.
-
Non-metals are more easily reduced.
-
They are termed “oxidizing agents.”
-
If a substance is reduced, its oxidation number decreases.
B. Oxidation Numbers
Rules for Assigning Oxidation Numbers:
(1)
Oxidation numbers for atoms that are free elements are always _________.
(2)
The oxidation number of monatomic ions are the same as the charge on the __________.
(3)
The sum of the oxidation numbers in a compound is always __________.
(4)
The sum of the oxidation numbers in a polyatomic ion is equal to ________________.
(5)
The oxidation number of Group 1 metals is always __________.
(6)
The oxidation number of Group 2 metals is always __________.
(7)
Oxygen almost always has a ______ oxidation state.* Unless in a peroxide
(8)
Halogens usually have a ______ oxidation state.*
(9)
Hydrogen usually has a ______ oxidation state.*
* There are rare exceptions to these rules.
Practice
Assign oxidation numbers to each element in the following compounds or polyatomic ions. To really keep on your toes, see if you
can name them as well.
1. MgBr2
9. CuSO4
17. K2Cr2O7
2. Cu
10. Cr
18. Al2O3
3. Fe2O3
11. H2CO3
19. Fe(NO3)3
4. AlN
12. Ba(NO3)2
20. SrCO3
5. SO3
13. NF3
21. Na2SO3
6. PO43-
14. CO
22. Ca(ClO)2
7. Cr2O72-
15. CO2
23. H2O
8. HClO2
16. CH4
24. K2S
C. Identifying REDOX Reactions
Identifying REDOX Reactions Using Oxidation Numbers
When a REDOX reaction occurs, there must be an element that is reduced (gained electrons) and an element that is oxidized (lost
electrons).
Are the following reactions REDOX reactions?
HINT: Find oxidation numbers of each element.
(1)
2H2 + O2  2H2O
(2)
HCl + NH3  NH4Cl
Oxidation Is Loss, Reduction Is Gain, “OIL RIG”
Identifying Oxidation and Reduction Half Reactions
Oxidation and reduction occur simultaneously. Elements and charges must be balanced for each half-reaction.
1.
Zn + CuCl2  Cu + ZnCl2
Oxidation Half:
Reduction Half:
2.
Na + Cl2  NaCl
Oxidation Half:
Reduction Half:
3.
Ca + O2  CaO
Oxidation Half:
Reduction Half:
4.
MnO2 + HCl  MnCl2 + H2O + Cl2
Oxidation Half:
Reduction Half:
5.
Cu + Ag+  Ag + Cu2+
Oxidation Half:
Reduction Half:
D. Balancing REDOX Reactions Using Half-Reactions
Steps to Balancing REDOX Reactions
When a REDOX reaction occurs, there must be an element that is reduced (gained electrons) and an element that is oxidized (lost
electrons). The number of electrons lost in the oxidation must equal the number of electrons gained in the reduction.
(1)
(2)
(3)
(4)
(5)
(6)
(7)
(8)
Assign oxidation numbers.
Decide what is being oxidized and what is being reduced.
Write the half reaction for each oxidation and reduction.
Balance all atoms
Balance the charge with electrons in each half reaction.
Balance the number of electrons gained with the number of electrons lost.
Add the two half reactions.
Simplify each equation by canceling out things that are the same on both sides.
Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons), “OIL RIG”
Oxidation Step: Oxidation number of element increases.
Reduction Step: Oxidation number of element decreases.
Hints:
1.
2.
There should be no electrons in your final equation.
The number of each element should be the same on both sides of the arrow when you are finished.
Balancing REDOX Reactions
Oxidation and reduction occur simultaneously. Elements and charges must be balanced for each half-reaction. Electrons must be
the same on either sides of the arrow.
1.
Zn + CuCl2  Cu + ZnCl2
Oxidizing Agent
Reducing Agent
Oxidizing Agent
Reducing Agent
Oxidation Half:
Reduction Half:
Balanced Reaction:
2.
Ni + Sn4+  Ni2+ + Sn
Oxidation Half:
Reduction Half:
Balanced Reaction:
3.
Ni + Sn+4  Ni+2 + Sn
Oxidizing Agent
Reducing Agent
Oxidizing Agent
Reducing Agent
Oxidizing Agent
Reducing Agent
Oxidizing Agent
Reducing Agent
Oxidizing Agent
Reducing Agent
Oxidation Half:
Reduction Half:
Balanced Reaction:
4.
Hg + Ag+1  Ag + Hg+2
Oxidation Half:
Reduction Half:
Balanced Reaction:
5.
KClO3  KCl + O2
Oxidation Half:
Reduction Half:
Balanced Reaction:
6.
H2S + O2  SO2 + H2O
Oxidation Half:
Reduction Half:
Balanced Reaction:
7.
H+ + Sn  Sn+2 + H2
Oxidation Half:
Reduction Half:
Balanced Reaction:
Electrochemistry
Page |6
E. Electrochemical Cells
The Basics of an Electrochemical Cell
Every electrochemical cell (contains two compartments where REDOX reactions can occur) contains 2 electrodes. An electrode is a
strip of metal that facilitates the loss or gain of electrons. There are two types of electrodes, known as the cathode and the anode.
Anode
Cathode
Where ___________________ occurs.
Where ___________________ occurs.
Electrons always flow from the ___________________ to the ___________________.
If you have 2 electrodes made of different metals connected, how can you tell which one is oxidized and which one is reduced? Table J!
Oxidation
If the metal is ______________ on Table J. (More active = easier to lose electrons)
Reduction
If the metal is ______________ on Table J. (Less active = harder to lose electrons)
Voltaic or Galvanic Electrochemical Cells
In a Galvanic/Voltaic cell, reactions occur ________________________; therefore, no additional energy must be added.
Spontaneous Reaction
- Reactions that proceed without the addition of energy.
- Predicting Spontaneous Reactions (Table J)
- If the metal by itself if higher up on the chart than the ion in the compound.
Electrochemistry
Page |7
-
In a voltaic cell, the anode is denoted with a NEGATIVE sign.
In a voltaic cell, the cathode is denoted with a POSTIVE sign.
-
Each half cell is connected by a salt bridge.
o What is a salt bridge?
-
Negative ions always flow from the _____________________________ to the _________________________.
-
Positive ions always flow from the _____________________________ to the _________________________.
-
The electrodes are connected by a metal wire which allows for the flow/movement of electrons. Remember: Electrons
always flow from the anode to the cathode.
-
A voltmeter usually measures the electron potential in Volts.
Electrochemistry
Page |8
Determining Spontaneity
Which of the following reactions will take place spontaneously?
a. Ni (SO4) + Pb 
b. Sr(CO3) + Sn 
Ni + Pb(SO4)
Sr + Sn(CO3)
c. Au (PO4) + Al 
Au + Al (PO4)
d. Fe(OH) + Cu 
Fe + Cu(OH)
Labeling Electrochemical Cell Processes
Between each of the following electrodes:
Species Oxidized
Calcium and
Iron
Silver and Nickel
Magnesium and
Lead
Copper and
Silver
Species Reduced
Anode Metal
With Sign
Cathode Metal
With Sign
Direction of eFlow
Direction of
Cation Flow in
the Salt Bridge
Electrochemistry
Page |9
Electrolytic Cells
Reactions that occur in an electrolytic cell occur ____________________________________________. You must force the reaction
to occur by adding electricity.
In Electrolytic Cells:
-
The anode is POSITIVE.
-
The cathode is NEGATIVE.
____________________
____________________
Why would you ever want to force a nonspontaneous reaction?
(a) To obtain pure metals
Many metals are only found as compounds in nature. Electrolysis can lead to a deposit of the pure metal on the cathode.
(b) To recharge a battery
A car battery powers the car through a spontaneous reaction, but what can you do if the battery dies?
(c) To coat one metal on top of another one, as with jewelry, or exhaust pipes.
a. To make something look more expensive or shinier
b. To improve corrosion resistance
Electrochemistry
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Using Electrolysis for Electroplating
An example of electrolytic cells in which an electric current is used to plate an object.
The substance to be plated is at the _____________________________________.
The substance that does the plating is at the ______________________________.
If I wanted to gold plate a necklace, where should I put the gold and where should I put the necklace in the electrolytic cell?
Why is it necessary to use an electrolytic cell to gold plate almost all other metals instead of a voltaic cell?
Would I need to use an electrolytic cell to plate nickel on top of an iron pipe? How do you know?
Electrochemistry
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Labeling Electrolytic Cells
Label the following for each of the electrolytic cells below.
a. Anode
b. Cathode
c. Reduction Electrode
d. Oxidation Electrode
e. Show electron flow using arrows.
f. Label ions.
g. Has the anode increased or decreased in mass?
h. Has the cathode increased or decreased in mass?
(a) Zn/Cu in CuSO4
(b) Ni/Ag in AgNO3
(c)
(d) Ag and a ring in AgNO3
Cu/Pb in CuCl2
Electrochemistry
P a g e | 12
Calculating Cell Potentials
It is possible to calculate the potential difference between two half-cells using standard reduction potentials. All standard reduction
potentials are tabulated relative to the Standard Hydrogen Electrode (SHE). The more positive the reduction potential value, the
more likely the substance is to undergo reduction.
Electrochemical Potential or Electromotive Force (emf)
E°cell = E°red (cathode) - E°red (anode)
Determining Cell Potentials
Predict the voltages or emf produced by the following cells.
1.
Zn/Zn+2//Fe+2/Fe
2.
Mn/Mn+2//Br2/Br-
Electrochemistry
3.
Ni/Ni+2//Hg2+2/Hg
4.
Cu/Cu+2//Ag+/Ag
5.
Using standard reduction potentials, calculate the standard emf for each of the following reactions:
P a g e | 13
(a) H2 (g) + I2 (s)  2H+ + 2I-
(b) Ni(s) + 2Ce+4 (aq)  Ni+2 (aq) + 2Ce3+(aq)
(c) Cr(s) + 2Cr+3 (aq)  3cr+2 (aq)
(d) 2Al3(aq)- + 3Cd(s)  2Al(s) + 3Cd2+(aq)]
6.
A voltaic cell is constructed: one electrode compartment has an aluminum strip in contact with a solution of Al(NO3)3, and
the other is a standard hydrogen electrode.
(a) Write the half reactions involved and determine which electrode is the anode and which is the cathode.
(b) Will the aluminum strip gain or lose mass as the cell operates?
(c) Write a balanced equation for the overall cell reaction.
(d) What is the standard emf of the cell?